Journal of The Electrochemical Society, 163 (10) A2377-A2384 (2016)
A2377
Rotating Ring-Disc Electrode Investigation of the Aprotic
Superoxide Radical Electrochemistry on Multi-Crystalline
Surfaces and Correlation with Density Functional Theory
Modeling: Implications for Lithium-Air Cells
Shrihari Sankarasubramanian,a,∗,z Jeongwook Seo,a Fuminori Mizuno,b,∗∗,c
Nikhilendra Singh,b and Jai Prakasha,∗∗
a Center
for Electrochemical Science and Engineering, Department of Chemical and Biological Engineering,
Illinois Institute of Technology, Chicago, Illinois 60616, USA
Research Institute of North America, Ann Arbor, Michigan 48105, USA
b Toyota
The realization of a practical Lithium-air cell depends on understanding the oxygen reduction reaction (ORR), identifying a stable
electrolyte and finding a suitable cathode. This study investigates the superoxide radical, its side reactions and correlates density
functional theory (DFT) predictions of the surface activity to the experimental kinetics. The ORR on glassy carbon (GC), multicrystalline Pt and multi-crystalline Au substrates in oxygen saturated 0.1 M Tetraethylammonium Perchlorate (TEAClO4 ) in Dimethyl
Sulfoxide (DMSO) was studied using cyclic voltammetery and the rotating ring-disk electrode (RRDE) technique. The RRDE data
was analyzed using kinetic models to understand the electrochemistry of the superoxide radical and calculate the surface reaction
rate constants. The percentage of the superoxide radicals detected at the ring from GC, Pt and Au surfaces correlated linearly to the
modeled activities of the substrates. Further, the modeled activity trend was found to correlate strongly with the ORR onset potential.
This study validates the DFT catalyst screening approach. It also shows that side reactions involving the superoxide radical occurs
in fresh, anhydrous DMSO based electrolytes.
© The Author(s) 2016. Published by ECS. This is an open access article distributed under the terms of the Creative Commons
Attribution 4.0 License (CC BY, http://creativecommons.org/licenses/by/4.0/), which permits unrestricted reuse of the work in any
medium, provided the original work is properly cited. [DOI: 10.1149/2.11161610jes] All rights reserved.
Manuscript submitted May 31, 2016; revised manuscript received July 22, 2016. Published September 1, 2016. This was Paper 230
from the San Diego, California, Meeting of the Society, May 29–June 2, 2016.
The Lithium-air cell is an attractive technology for beyond Li-ion
applications due to its high theoretical specific energy of 3505 Wh
Kg−1 which is significantly higher than current lithium-ion cells. To
achieve practical application as a replacement for Li-ion cells, current
Li-air cells must overcome problems of (i) poor cycle life; (ii) high
charge and discharge overpotentials resulting in low columbic efficiency and low power density;1 (iii) eventual ambient air operation.
The poor cycle life reported in many articles can be traced to the oxidative instability of the electrolyte2–4 in conjunction with electrode
passivation due to irreversible product deposits and high overpotentials for decomposing Li2 O2 .5,6
The issue of electrolyte instability has been extensively reported
since the early attempts to fabricate a secondary lithium-air cell adopting typical lithium-ion cell electrolytes. Alkyl carbonate based electrolytes were found to undergo nucleophilic attack, form passivating
films and consume the electrolyte.2 Extensive studies on a wide variety of electrolytes such as acetonitrile,7 ethers,3 dimethylformamide,4
dimethylsulfoxide (DMSO),8 and various ionic liquids9,10 have been
carried out and relatively stable electrolytes were identified.11 Further, the donor number of the electrolyte solvent was found to
be critical to the mechanism of the oxygen reduction reaction
(ORR).12 Thus 1,2 dimethoxyethane (DME) and dimethylsulfoxide
(DMSO) were extensively adopted for further study due to their high
donor numbers and relative stability. Ionic liquids were also studied due to similarly favorable properties and their inherently good
conductivity.9,10,13,14
Nevertheless, oxidative stability continues to be of concern. Recently, the stability of DMSO has been called into question. Sharon
et al. proposed multiple schemes for the reaction of DMSO with reactive oxygen species and Li2 O2 .15 Similar reactivity of DMSO with
Li2 O2 resulting in LiOH formation has also been reported taking electrolyte aging into account16 and by direct exposure of Li2 O2 to DMSO
and its subsequent characterization by X-ray photo-adsorption spec∗ Electrochemical Society Student Member.
∗∗ Electrochemical Society Member.
c
Present address: Battery Material Engineering & Research Division, Toyota Motor
Corporation, Higashifuji Technical Center, Shizuoka 410–1193, Japan.
z
E-mail: ssanka11@hawk.iit.edu
troscopy (XPS).17 While the evidence strongly suggests that DMSO
is indeed oxidatively unstable, water contamination could also be a
factor due to the well-known solubility of water in DMSO.18 This
would be especially pertinent as the scanning electron microscope
(SEM) images of Li2 O2 reported in Kwabi et al.16 showed nucleation
and particle structure whose occurrence has been shown to require the
presence of water19 and due to the electrolyte aging involved. Further,
the presence of water has been shown to result in LiOH being the
reduction product.20
The present study sought to examine the oxygen reduction reaction
(ORR) and identify any side reactions on glassy carbon (GC), multicrystalline Pt and multi-crystalline Au electrodes in anhydrous 0.1 M
TEAClO4 in DMSO using the Rotating Ring-Disk Electrode (RRDE)
technique. The absence of lithium ensured that the electrode surface
was pristine and allowed for the application of well-established kinetic
models21,22 to calculate the rate constants for reactions and to identify possible side and parallel reactions.23 These techniques have been
previously used by us for kinetic investigations of the H+ ORR on
various surfaces such as Lead Ruthenate pyrochlores,24 Ruthenium,25
and Cobalt-Palladium.26 The measurement of the ratio of superoxide
species in the solvent is challenging in the Li salt containing electrolyte due to the short half-life of the radical27 and the change in
the electrode surface due to product deposition. But these constraints
were overcome in the present study by the use of an unreactive TEA+
cation. This also allowed for the measurement of the degree of side
reactions between the superoxide radical and the electrolyte and show
that pristine, fresh DMSO also undergoes reactions with the superoxide radical. The ratio of superoxide radicals that are in the electrolyte
was calculated and correlated with the density functional theory (DFT)
modeling based activity of the substrates. The use of DFT for catalysts screening is well established28 but is also beset with challenges
in the Li-air system due to the requirement to model the solvent effect
and the changing substrate. Herein, we correlated the DFT predictions reported by us previously29 to the first electron transfer reaction
resulting in the formation of superoxide radical, which subsequently
reacts to form the peroxide product in the presence of Li+ via a variety of routes and possible side products due to nucleophilic attack of
some electrolyte solvents. Similar to our work on the ORR in polymer
electrolyte membrane fuel cells (PEMFCs),30 we found that relatively
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A2378
Journal of The Electrochemical Society, 163 (10) A2377-A2384 (2016)
simple DFT models showed a remarkable degree of correlation with
experimental data.
Experimental
The 0.1 M solution of Tetraethylammonium perchlorate
(TEAClO4 ) (Alfa Aesar) in Dimethylsulfoxide (DMSO) (>99.9%,
Fisher) was prepared in a MBraun argon filled glove box with H2 O
and O2 levels <0.5 ppm. The TEAClO4 salt was dried in a vacuum
oven for 24 hours at 40◦ C and DMSO was distilled, argon purged
and stored over zeolite beads in an argon filled glove box to ensure
absence of moisture before use. The DMSO was opened, stored and
handled in an Ar glove box with H2 O levels <0.5 ppm and data from
the literature shows that the rate of water uptake at room temperature
and in the absence of stirring was very low (0.05 wt% water uptake
after 1-hour exposure to air at 75% relative humidity (295 K)).41 In
the extreme case of exposure to 98% relative humidity (295 K) air
for 5 days, the mole fraction of water was 2.17.41,42 Thus, in light of
the precautions taken, the water content is expected to be close to the
rated value of less than 0.01%. Despite these precautions, absorption
of a small amount of water is possible and its impact on the reaction is
discussed in light of the “catalytic” effect reported in the literature.44
The electrochemical measurements were carried out in a setup that
consisted of a three-neck electrochemical cell with Teflon stoppers
with openings for working, counter and reference electrodes and a
gas purge line. The working electrode was a Pine instruments RRDE
assembly with a Pt ring and interchangeable disks of glassy carbon
(GC), Au and Pt polished to a mirror finish with 0.05 μm alumina
suspension. The counter electrode consisted of a coiled Pt wire in
a fritted glass tube (Pine instruments) and the reference electrode
consisted of an Ag wire dipped in a solution of 0.1 M AgNO3 in
CH3 CN (−3.85 V vs Li/Li+ in 0.1 M TEAClO4 /DMSO). A gas purged
bearing assembly (Pine instruments) with continuous Argon purge
was used to ensure isolation from the environment during RRDE
rotation. The entire setup was assembled in the glove box, sealed
and then transferred to a glove-bag which was kept at a positive
pressure relative to the atmosphere using a continuous argon feed
to prevent moisture contamination. The electrochemical impedance
spectra (EIS) were recorded using a Solartron analytical frequency
response analyzer to measure the resistance of the setup prior to
electrochemical measurements and was later used in iR correction of
the RRDE voltammograms. The electrochemical measurements were
performed using a multi-channel potentiostat (Solartron analytical).
Results and Discussion
Fig. 1 depicts possible pathways for the reduction of oxygen in
a lithium ion containing electrolyte. There were no reports in the
literature of Li2 O being detected in Li-O2 systems and so this product
of a possible 4-electron reduction was not considered in our analysis.
The initial step in the ORR mechanism was considered to be the
adsorption and subsequent reduction of oxygen to the superoxide
radical. Outer sphere electron transfer31,32 to form O2 − and the possible
subsequent adsorption of the superoxide radical was also considered.
This mechanism for forming the superoxide is a possibility in light
of the degradation of the electrolyte which could be caused by the
highly nucleophilic superoxide radical remaining in the electrolyte.43
Laoire et al.12 show that the stability and half-life of the superoxide
radical is a function of the donor number of the solvent and propose a
complex formation between the tetraalkylammonium cations and the
superoxide radicals. Due to the consequent stability of the superoxide
radical in high donor number solvents such as DMSO, its half-life
could be expected to increase greatly. The electrochemical detection
of the superoxide radical was expected to be a function of this degree
of stabilization. As depicted in the figure, with the stable superoxide
radical, surface adsorption would be entirely dependent on the binding
energy of the intermediate and would imply that disproportionation
reaction on both the surface35 and the electrolyte8,13 are possibilities.
Another possibility considered was the hybrid mechanism proposed
elsewhere.34 Li2 O2 has been observed on the surface of the cathode5
and has also been detected in the electrolyte.33 Based on observations
with H+ ORR in aqueous systems,24 it is reasonable to expect a
combination of these mechanisms occurring. The present system of
TEAClO4 /DMSO was selected so as to have a simple system for
the study of superoxide radical, its side reactions and the proportion
of reduced species that is desorbed into the electrolyte. Further, the
effect of the electrode material on the oxygen reduction reaction in
various aprotic electrolytes was found to be reported multiple times
in the literature through both experimental1,37,39,45,46 and theoretical
methods29,47 and one part of the present study serves to examine the
same in the context of the Li-O2 cell.
Cyclic voltammetry.—Cyclic voltammetry allowed the identification of the potentials at which reactions occur in this model system
under cathodic and anodic conditions. Parameters such as the peak
voltage, reaction onset potential, the separation between the reduction and oxidation peaks allowed for the initial characterization of
the system. Fig. 2 shows the cyclic voltammograms (CV) of GC, Pt
and Au discs in oxygen saturated 0.1 M TEAClO4 /DMSO at scan
rates of 50 mV s−1 , 100 mV s−1 and 500 mV s−1 . It was evident
from the voltammograms in the Ar saturated condition, in case of
all three substrates, that the electrolyte was electrochemically stable
within the potential window of interest. Outside this window, oxidation peaks were observed between 3.7 to 4 V as clearly seen in the slow
5 mV s−1 scan CVs depicted in Fig. 3b. This was similar to the potential at which Sawyer et al.39 reported oxidation peaks in the presence
of protons (from added HClO4 ) and also close to the potential at which
oxidation peaks were observed with the addition of water.44 The voltage range from 1.85 V to 4.35 V covered the operating range of Li-air
cells wherein the discharge and charge plateaus were found to occur
Figure 1. Reaction scheme for the Oxygen reduction reaction in a Li-air cell. The hypothesized disproportionation reactions on the surface (black, solid arrows)
and the electrolyte bulk (red, dot-dash arrows) is depicted.
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Journal of The Electrochemical Society, 163 (10) A2377-A2384 (2016)
Figure 2. Cyclic Voltammograms of O2 saturated 0.1 M TEAClO4 in DMSO
on Glassy Carbon (GC), Pt and Au working electrodes. The scan rates are:
50 mVs−1 (dashed line), 100 mVs−1 (dotted line), 500 mVs−1 (solid line).
The background current in a Ar-saturated condition is shown as a solid line.
at ∼2.5 V and ∼4 V respectively. Upon saturation with oxygen, the
voltammograms showed a reduction peak below 2.45 V and oxidation
peak above 2.65 V during the cathodic and anodic scans respectively
on GC, Pt and Au, with a voltage separation whose magnitude suggested a degree of irreversibility. The shift in the peak position with
the scan rate was found to be in agreement with the proposal of Bard
and Faulkner36 of a 30/α mV shift for a 10-fold increase in scan rate.
Thus, the reduction and oxidation peaks were believed to correspond
to the reduction of adsorbed oxygen to the superoxide radical and its
subsequent oxidation to oxygen respectively. The Pt disc was found to
show the highest onset potential for the ORR at 2.9 V during the cathodic scan while the Au disc showed the lowest onset potential for the
oxygen evolution reaction (OER) at 2.5 V for the anodic scan. Thus,
the advantages of a Pt-Au bifunctional catalyst37 was readily apparent.
Rotating ring-disc electrode (RRDE) measurements.—The rotating ring-disc electrode allowed for the characterization of the
adsorption-desorption of intermediates on the disc and also allowed
for the identification of the pathways of the reduction reaction occurring on the disc using a kinetic model. The ring, which was held
at an oxidizing potential as identified from the CVs, allowed for the
possible identification of various intermediates. The RRDE measurements depicted in Fig. 3a were carried out on GC, Pt and Au discs
respectively scanning at 5 mV s−1 between 1.85 V to 2.85 V while the
Pt ring was held at 3.35 V. The scan rate allowed sufficient time for
double layer relaxation and also served to mitigate experimental noise.
The disc voltage window was chosen so as to have the linear sweep
voltammogram (LSV) cover the entire range of the reduction reaction
while the ring was held at an anodic potential to ensure complete
oxidation of superoxide radicals and any other side reaction products
A2379
produced. The disc currents showed very little hysteresis except in
case of the platinum disc. This was noted as being in line with common observations in the fuel cell literature wherein the reduction of
oxide impurities on the platinum surface in the cathodic scan results in
higher currents on the anodic scan. The relation
√ between the rotation
rate and the limiting current, where iL ∝ ω (eg: iL@1600RPM = 2.
iL@400RPM ), was indicative of a convective-diffusive regime. Further,
since the above proportional relation is dependent on the reaction surface being unchanged, it was seen to be evidence of the surface not
being affected by deposition of any possible side-reaction products.
The transient time between the disc and ring was found to be 6 s,
12 s and 34 s for the GC, Au and Pt discs respectively. The transient
time was taken to be the time between the onset of the ORR on the
disc and the increase of the ring current above the background current
during the cathodic scan. The cathodic scan was chosen because the
rate of transport of the reduced species to the ring was thought to be
of a different order as compared to the rate of the OER on the ring.
Thus, during the cathodic scan the ring was expected to respond as
soon as reduced species reached the ring, In the anodic scan it was
expected that the ring current would persist till all the reduced species
at the ring has been oxidized despite there being no transport of new
reduction products. Thus, measuring the transient time during the
cathodic scan was thought to be most accurate. Since the electrolyte,
the convection currents (a function of the RRDE rotation rate) and the
potential gradient between the ring and the disc were the same in all
three cases, the difference in the transient times was hypothesized to be
purely a function of the concentration gradient of the reduced species
between the near-disc environment and the ring. The creation of such a
gradient would require that the product desorption rates be different for
the different substrates in line with variation of intermediate binding
energies. The percentage of product species transported to the ring
was investigated to test this hypothesis.
The RRDE electrode was rated to have a collection efficiency (N)
of 0.25 and this can be used to calculate the proportion of the ions
produced on the disc that desorb and travel to the ring using the
following equation:
χ (%) =
i Ring
N
i disk +
i Ring
N
∗ 100
[1]
The percentage of the ions produced on the disc diffusing away
from the GC, Au and Pt disks were found to be 29%, 27% and 22%
respectively. Thus it was found that the transient times for the various
substrates were inversely proportional to the percentage of reduced
species that desorbs. The case of outer sphere electron transfer was
also considered where the superoxide radical formed does not adsorb
on the disc but undergoes side reactions in the near disc environment.
Thus, the superoxide radical detected at the ring was thought to be
the unreacted superoxide. The χ (%) data showed that 88%, 83% and
81% of the superoxide radicals produced on or near the Pt, Au and
GC electrodes respectively reacted further. This was found to follow
the same reactivity trend as reported elsewhere.43 Thus, the analysis
presented herein adequately explains the cases of both electron transfer
to adsorbed oxygen on the electrode surface and also possible outersphere electron transfer. Further analysis of the impact of the disc
material on the adsorption of the reaction intermediates using binding
energies calculated from DFT is discussed in section Electrochemical
reaction kinetics.
Koutecky-Levich plots.—The characteristic LSVs obtained using
the rotating ring-disc electrode were well described by the classic
Koutecky-Levich equation. The Koutecky-Levich equation describes
the disc current of an RRDE system if the reaction kinetics is firstorder with respect to dissolved oxygen. In that case, the observed
currents at the rotating disk electrode are related to the rotation rate ω
by the following equation:
1
1
1
1
1
= +
= + √
i
ik
iL
ik
B ω
[2]
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A2380
Journal of The Electrochemical Society, 163 (10) A2377-A2384 (2016)
Figure 3. (a) Rotating ring-disk voltammograms on (i) Glassy carbon (GC), (ii) Pt and (iii) Au disks recorded at a scan rate of 5 mVs−1 in O2 saturated
0.1 M TEAClO4 in DMSO. Ring currents recorded with the Pt ring held at 3.35 V vs Li/Li+ . The solid lines and dashed lines indicate cathodic and anodic scans
respectively; (b) Cyclic Voltammograms of O2 saturated 0.1 M TEAClO4 in DMSO on (i) Glassy carbon (GC), (ii) Pt and (iii) Au disks at a scan rate of 5 mVs−1 .
where ik is the kinetic current (i.e., the maximum current attainable in
the absence of any mass-transport limitation), iL is the mass transport
limiting current and B is a constant. They are given by:
i k = (i L . i)/ (i L − i)
2
1
B = 0.62 n F AD 3 υ− 6 Cb
−1
[3]
[4]
where F is 96,485 C mol , n is the number of electrons in the overall
reaction, A is the disc area, D is the diffusion coefficient of O2 , Cb is
the bulk concentration of O2 , υ is the kinematic viscosity, and ω is
the rotation
rate in revolutions per minute (rpm). Thus a plot of 1/i
√
vs. 1/ ω for various potentials would be expected to yield straight
lines with intercepts corresponding to ik and the slopes yielding the B
values. Knowledge of D, Cb and υ for a given electrolyte along with
the disc area enables the calculation of the number of electrons. In the
present system, D = 3.11 × 10−5 cm2 s−1 ,38 Cb = 2.1 mM cm−339 and
υ = 9.57 × 10−6 cm2 s−1 .40
The disc currents showed three distinct behaviors across the voltage window for reduction. Initially, the currents were independent
of the rotation rate in a kinetic limited regime where the iL term in
Equation 2 was zero. At more cathodic potentials, the currents en-
tered a mixed-control regime where they began to display a rotational
dependence. Finally, in the
√ mass-transport controlled regime, the currents displayed the iL ∝ ω relationship characteristic with the 1/ik
intercept approaching zero. The mixed control regime was the region
of interest for calculating the number of electrons using Equation 4.
The Koutecky-Levich plots are shown in Fig. 4i. In the mixed control
regime between 2.3 V and 2.7 V, the plots were seen to be almost
parallel and showed close to one electron transfer within the experimental error of 10%. At 2.3 V the intercept was seen to be almost
zero, indicative of the beginning of the mass-transport limited regime.
In the kinetic controlled regime, the rate constant of the disc reaction
maybe calculated from the intercept of the K-L plots using the relation i = n F AkCb .36 The rate constants calculated by this method are
tabulated in Table I.
Tafel plots.—The Tafel equation was used to elucidate the rate
determining step in the multi-step reduction reaction occurring on the
disc. The Tafel equation relates the kinetic current and the potential
in the kinetic limited regime of the surface electrochemical reaction.
It is as follows:
η = a + b.log (i k )
[5]
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Journal of The Electrochemical Society, 163 (10) A2377-A2384 (2016)
A2381
Figure 4. (i) Koutecky-Levich plots for oxygen reduction on (a) glassy carbon, (b) Pt, (c) Au in O2 saturated 0.1 M TEAClO4 / DMSO solution during the anodic
scan. The error bars show a 5% standard error.; (ii) Representative Tafel plots for oxygen reduction at 900 RPM on (a) glassy carbon, (b) Pt, (c) Au in O2 saturated
0.1 M TEAClO4 / DMSO solution during the anodic scan.
Where,
a=
2.3RT
.logi 0
α.F
[6]
2.3RT
α.F
[7]
b=−
Table I. Representative comparison of Oxygen reduction reaction
rate constants calculated by the Kinetic model presented herein
and the intercepts of the classical Koutecky-Levich (K-L) plot, in
the kinetic region (2.7 V vs Li/Li+ ). The oxygen concentration was
taken to be Cb = 2.1 mM cm−3 .39
Substrate
iK
(mA cm−2 )
K (cm s−1 )
(from K-L plot)
K (cm s−1 )
(from kinetic model)
GC
Pt
Au
0.213
0.568
0.447
1.05 × 10−3
2.80 × 10−3
2.21 × 10−3
1.29 × 10−3
3.25 × 10−3
1.70 × 10−3
Here, α is the transfer coefficient and i0 is the exchange current density
in mA cm−2 . The variable b in Equation 7 is the Tafel slope.
Fig. 4ii depicts the mass-transport corrected Tafel plots constructed
from the disc current data shown in Fig. 3a. The kinetic currents were
obtained from Equation 3. All three disc materials showed a Tafel
slope of 120 mV dec−1 which was indicative of the first electron
transfer step being rate determining, as would be expected for the
formation of the superoxide radical. The 120 mV dec−1 Tafel slope
was observed for over two orders of magnitude of the kinetic current
which was indicative of true Tafel behavior.
N(iD /iR ) vs ω−1/2 plots.—Damjanovic et al.23 provide one of the
first, comprehensive discussion of the use of N(iD /iR ) vs ω−1/2 plots to
study the side reactions using a RRDE. This treatment was later used
in the works of Wroblowa et al.22 and Hsueh et al.21 to study the ORR
in aqueous systems.
The case of an ideal system with no side reactions or loss of
products was considered where the ring current would be equal to the
collection efficiency of the system times the disc current. This implied
that, subject to the physical constraint of the collection efficiency of
the system, all the reduction products would be oxidized at the ring.
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A2382
Journal of The Electrochemical Society, 163 (10) A2377-A2384 (2016)
Thus, a plot of N(iD /iR ) vs ω−1/2 would result in a single straight line
and the plots at various potentials would overlap. The occurrence of
side or parallel reactions and the deposition of the product on the disc
or its precipitation would result in deviation from the ideal described
previously.
The second case examined was that of disc products that react further. In the absence of convection (no rotation) the ring current would
depend on the rate of product desorption and diffusion to the ring
which would be independent of the potential and hence the N(iD /iR )
vs ω−1/2 plots converge to a single intercept value at ω−1/2 = 0. It was
seen that as the rotation rate increased, the movement of product to
the ring was no longer transport limited and hence the N(iD /iR ) plots
diverged from their common intercept.
In the case of parallel reactions where the products do not react
further, the net rate of reaction was thought to be potential dependent as
each reaction would be a function of potential. But as the products do
not react further, the concentration in the near electrode environment
would be only a function of the rate of desorption and transport to
the ring. Thus the N(iD /iR ) vs ω−1/2 plots would be a series of straight
lines parallel to the x-axis and the intercept would be a function of
potential.
A fourth case analyzed was that of parallel reactions producing
products that can react further. Similar arguments as for the third case
were found to apply. But now the N(iD /iR ) vs ω−1/2 plots were expected
to show a dependency on the convection rate due to the competition
between further reaction of the product and its transport to the ring.
The N(iD /iR ) vs ω−1/2 plots shown in Fig. 5 were found to be of
the fourth case. The plots values were taken from the mass transport
limited region so as ensure the transport and kinetics are at steady
state. These plots indicated that the superoxide radical does undergo
further reaction. The reaction was thought to consist of nucleophilic
reaction with DMSO and reaction with any traces of water present.
Further studies with techniques such as in-situ XPS are required to
identify the exact products produced.
iL /(iL -iD ) vs. ω−1/2 plots.—Following the work of Wroblowa
et al.22 which was extended by Hsueh et al.21 the rate constants for
the formation of the superoxide radical were calculated. The equations given in the above studies were adopted for the present case as
follows:
k1
iL
1
= 1 + ω− 2
i L − id
Z
2
[8]
−1
6
Where Z = 0.62 D O3 2 ν D M
S O in which DO2 represents the diffusion
coefficient of oxygen and νDMSO is the viscosity.
Thus the slope of the iL /(iL -iD ) vs. ω−1/2 plots shown in Fig. 6 were
used to calculate the rate constants for superoxide formation. The
model predicted that the intercept was 1 and this was borne out in the
case of all the substrates. The model was applied in the kinetic region
of the disc voltammograms of Fig. 3a. Beyond the kinetic region the
model would be unphysical as it would track the disc current increase
and predict a constantly rising rate constant whereas the increase in
the disc current is attributable to the improved mass transport with the
kinetics reaching its steady state.
Electrochemical reaction kinetics.—The rate constants calculated
from the model described in section iL /(iL -iD ) vs. ω−1/2 plots are shown
in Fig. 7. These rate constants were compared with rates calculated
from the classical Koutecky-Levich equation and found to be in good
agreement as shown in Table I. The rate constants were seen to increase in the order GC < Au < Pt with the values at the start of the
mixed control region being 0.0013 cm s−1 , 0.0018 cm s−1 , 0.0034
cm s−1 respectively. The predicted first electron transfer activity of
the carbon basal plane, Au (111) plane and Pt (111) plane from DFT
calculations29 was compared to the experimental values and the trends
were found to be in close agreement. The experimental values were
for multi-crystalline electrodes but the electrochemical reactions tend
to be dominant on a single plane and thus the DFT models of the most
Figure 5. N(iD /iR ) vs ω−1/2 plots for oxygen reduction on (i) glassy carbon,
(ii) Pt, (iii) Au in O2 saturated 0.1 M TEAClO4 / DMSO solution. The error
bars show a 5% standard error.
active crystal plane was seen to be in close agreement with experiments. It was hypothesized in Sankarasubramanian et al.29 that since
the maximal surface activity is a function of the intermediate binding
energy, the binding energies could serve as a good predictor of surface
activity. It was seen from the calculated rate constants and the onset
potentials of the ORR shown in Fig. 8i that the predicted activities
of the substrates closely match experimental trends. Further, the relation between the activity, the ORR onset potential and the percentage
of disc species that desorbs was examined. The χ (%) served as an
indicator of the reduced species binding energy, the degree of side
reactions and it was found that the theoretical activity and the experimental onset potentials showed a linear relationship with χ within
experimental and modeling error of 10% as seen from Figs. 8ii and
8iii. Finally, the experimental data was found to closely fit the kinetic
model predictions which are predicated on the reaction substrate being
unchanging. Thus, it can be concluded that any side-reaction product
deposits had little or no effect on the kinetics measured herein.
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Journal of The Electrochemical Society, 163 (10) A2377-A2384 (2016)
Figure 6. iL /(iL -iD ) vs. ω−1/2 plots for oxygen reduction on (i) glassy carbon,
(ii) Pt, (iii) Au in O2 saturated 0.1 M TEAClO4 / DMSO solution. The error
bars show a 10% standard error.
A2383
Figure 8. (i) Representative ORR polarization curves on GC, Pt and Au in O2
saturated 0.1 M TEAClO4 in DMSO at a scan rate of 5 mVs−1 and RDE rotation
rate of 1600 RPM; (ii) Correlation between surface activity and percentages of
superoxide radicals detected at ring and (iii) Correlation between ORR onset
potential and percentages of superoxide radicals detected at ring. The error
bars show a 10% standard error. The DFT model surface activity values were
taken from Reference 29.
Conclusions
Figure 7. Rate constants for the one electron oxygen reduction reaction to
form the superoxide radical in the kinetic controlled region. The values are
calculated from the slopes of the iL /(iL -iD ) vs. ω−1/2 plots depicted in Figure 8.
The error bars show a 10% standard error.
The kinetics of the oxygen reduction reaction in 0.1 M
TEAClO4 /DMSO was studied using Rotating Ring-Disk voltammetry.
The cyclic voltammetry peaks for reduction of oxygen and subsequent
oxidation of the superoxide radical was found to show a voltage difference that is possibly indicative of irreversibility and in agreement with
reports in the literature that observe high overpotentials for charge and
discharge. Only a portion of the superoxide radicals produced on or
near the disc was detected at the ring indicating that both further electrochemical and chemical reduction of the superoxide radicals in the
presence of Li+ or other alkali metal ions could be expected. This was
also an indicator of the occurrence of possible nucleophilic reactions
between the superoxide radical and the electrolyte. The disc voltammograms were analyzed using standard kinetic models and it was
seen that the superoxide radical undergoes further reaction, possibly
with the DMSO solvent. The rate constants for the formation of the
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A2384
Journal of The Electrochemical Society, 163 (10) A2377-A2384 (2016)
superoxide radical were calculated and found to be in close agreement with prior predictions for the first electron transfer form DFT
calculations. Further, it was seen that the hypothesis of the intermediate binding energy being a good predictor of substrate activity was
borne out and seems to point to a possible screening method for potential catalytic surfaces for preferential electrochemical reduction of
oxygen.
Acknowledgments
This work was supported by the Toyota Research Institute of North
America (TRINA). The authors thank Dr. Aadil Benmayza and Dr.
Javier Parrondo for fruitful discussions.
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